Bond energy is the amount of energy required to break one mole of a specific chemical bond in a gaseous state. It is a measure of bond strength in a chemical bond and typically expressed in kJ/mol.
In any chemical reaction bond breaking requires energy (endothermic), while bond formation releases energy (exothermic). The overall energy change depends on the balance of these two processes.
Enthalpy (ΔH) is the heat content of a system at constant pressure. It indicates the heat absorbed or released during a reaction. ΔH < 0 (heat is released) indicates Exothermic reaction and ΔH > 0 (heat is absorbed) indicates Endothermic reaction.
The enthalpy change of a reaction can be estimated using bond energies: ΔH = Total bond energies of bonds broken - Total bond energies of bonds formed.
The heat of reaction is the change in enthalpy during a chemical reaction. It represents the net energy change when reactants are converted to products.
Bond energy is the average energy needed to break a bond in a molecule, while bond dissociation energy (BDE) refers to the exact energy required to break a specific bond in a particular molecule.
Bond energy depends on factors such as bond length, bond order, atomic size, and electronegativity.
Using a calorimeter, the heat change (q) is measured, often using the formula: q = m · c · ΔT.
Double and triple bonds involve more shared electrons, which increases the bond strength and decreases bond length.
An exothermic reaction releases energy to the surroundings, making the environment warmer.
An endothermic reaction absorbs energy from the surroundings, making the environment cooler.
If bond formation releases more energy than bond breaking requires, the reaction is likely spontaneous. However, other factors like entropy must also be considered.
Hess’s Law states that the total enthalpy change of a reaction is the same, regardless of the path taken.